Acids and bases may be strong or weak. This classification, however, is not tied to the concentration of their solutions. For example, citric acid is weak enough as an acid such that we can drink solutions like fruit juicesthat contain high concentrations of the acid.
What Makes a Strong Acid or Strong Base?
Strong electrolytes are completely dissociated into ions in water. The acid or base molecule does not exist in aqueous solution, only ions. Weak electrolytes are incompletely dissociated.
Strong Acids
Strong acids completely dissociate in water, forming H+ and an anion. There are six strong acids. The others are considered to be weak acids. You should commit the strong acids to memory:
HCl - hydrochloric acid
HNO3 - nitric acid
H2SO4 - sulfuric acid
HBr - hydrobromic acid
HI - hydroiodic acid
HClO4 - perchloric acid
100% dissociation isn't true as solutions become more concentrated. If the acid is 100% dissociated in solutions of 1.0 M or less, it is called strong. Sulfuric acid is considered strong only in its first dissociation step.
H2SO4 -> H+ + HSO4-
Weak Acids
A weak acid only partially dissociates in water to give H+ and the anion. Examples of weak acids include hydrofluoric acid, HF, and acetic acid, CH3COOH. Weak acids include:
Molecules that contain an ionizable proton. A molecule wih a formula starting with H usually is an acid.
Organic acids containing one or more carboxyl group, -COOH. The H is ionizable.
Anions with an ionizable proton. (e.g., HSO4- --> H+ + SO42-)
Cations
transition metal cations
heavy metal cations with high charge
NH4+ dissociates into NH3 + H+
Strong Bases
Strong bases dissociate 100% into the cation and OH- (hydroxide ion). The hydroxides of the Group I and Group II metals usually are considered to be strong bases.
LiOH - lithium hydroxide
NaOH - sodium hydroxide
KOH - potassium hydroxide
RbOH - rubidium hydroxide
CsOH - cesium hydroxide
*Ca(OH)2 - calcium hydroxide
*Sr(OH)2 - strontium hydroxide
*Ba(OH)2 - barium hydroxide
* These bases completely dissociate in solutions of 0.01 M or less. The other bases make solutions of 1.0 M and are 100% dissociated at that concentration. There are other strong bases than those listed, but they are not often encountered.
Weak Bases
Examples of weak bases include ammonia, NH3, and diethylamine, (CH3CH2)2NH.
Most weak bases are anions of weak acids.
Weak bases do not furnish OH- ions by dissociation. Instead, they react with water to generate OH- ions.
Changes in the acidity or basicity of some solutions can be critical in some systems such as the following:
a body of water supporting aquatic plants and animals
our stomachs digesting ingested food
cells of living organisms where vital life processes are carried out
some industrial processes where reactions take place at a desired rate
As discussed earlier, the acidity of a solution is measured by the hydronium ion concentration. For most solutions this is a very small number, usually expressed using exponential notation. In 1909, Soren Sorensen, a Danish biochemist, proposed the pH scale as a more convenient way of expressing hydronium ion concentration. The symbol pH stands for some German words which literally mean ' the power of the hydrogen ion.'
The pH of a solution is defined as the negative of the logarithm of the hydronium ion concentration.
pH = -log [H3O+]
Pure water has a hydronium ion concentration of 1 x 10 raised to -7. Thus, using the formula given above, we can now say that pure water has a pH of 7.
Most cleaning solutions are acids and bases and there is usually a right cleaning solution for a particular stain. Reading the label of the solution will give you an idea which product is the best for the stain that you will remove. Therefore, we must be familiar with the formulas of acids and bases and their respective names.
Naming Acids
Acids like hydrogen chloride which are composed of only two elements, hydrogen, and a nonmetal, are called binary acids. The name hydrogen chloride refers to the pure compound which is a gas. When hydrogen chloride is dissolved in water, the solution formed hydrochloric acid. Other examples are the following:
Pure Acid
Aqueous Solution
HF(g)—hydrogen fluoride
HF(aq)—hydrofluoric acid
HI(g)—hydrogen iodide
HI(aq)—hydroiodic acid
Many other acids, called ternary acids, consist of three elements—hydrogen and two other nonmetals. Many of these acids are oxyacids, that is, one of the two nonmetals is oxygen. Examples are nitric acid and sulfurous acid.
The names of these oxyacids are based on the number of oxygen atoms per molecule.
Naming Oxyacids of Sulfur
Formula
Name
Formula and Name of Oxyanion Formed upon Reaction with Water
H2SO3
Sulfurous acid
SO3-2, sulfite
H2SO4
Sulfuric acid
SO4-2, sulfate
Take note of the progression of the change in the names of the acids in each set as related to the progression in the number of oxygen atoms per molecule of the acid.
Naming Oxyacids of Chlorine
Scheme for Naming the Acids
Formula
Name
Formula and Name of Oxyanion Formed upon Reaction with Water
Hypo___ous
HClO
Hypochlorous acid
ClO-, hypochlorite
___ous
HClO2
Chlorous acid
ClO2-, chlorite
___ic
HClO3
Chloric acid
ClO3-, chlorate
Per___ic
HClO4
Perchloric acid
ClO4-, perchlorate
Common Acids and Their Uses
Name
Uses
Acetic acid
Acid in vinegar used to season and preserve food; cleans and deodorizes
Hydrochloric acid
Aids in digestion; used as toilet bowl cleaner and for cleaning metal surfaces
Sulfuric acid
Used in automobile batteries and in making dyes and plastics; a dehydrating agent
Nitric acid
Used in making explosives and fertilizers
Phosphoric acid
Removes hard water deposits; used in making fertilizers; used in softdrinks (in dilute form)
Carbonic acid
Used in the manufacture of carbonated drinks
Acetylsalicylic acid (aspirin)
Reduces pain and inflammation
Naming Bases
Bases are compounds consisting of a metal and a hydroxide ion. They dissociate into these ions when dissolved in water. Examples of base dissociation reactions are:
NaOH H2O Na+(aq) + OH-(aq)
Ca(OH)2(s)H2O Ca2+(aq) + 2 OH-(aq)
Bases are named just like binary acids in pure form; the ions are named, then combined.
Example 1: NaOH
Ions: Na+- sodium
OH- - hydroxide
Base: sodium hydroxide
Example 2: NH4OH
Ions: NH4+- ammonium
OH- - hydroxide
Base: ammonium hydroxide
Common Bases and Their Uses
Name
Uses
Sodium hydroxide
Used in making soaps and detergents; a drain and oven cleaner
Lithium hydroxide
Used in removing CO2 from air in confined areas such as submarines and spaceships
Magnesium hydroxide (in Milk of Magnesia)
Used as an antacid in small dosages and laxative in large amounts
Aluminum hydroxide
Used as an antacid with no dosage restrictions
Ammonia
Used in the production of fertilizers and cleaning solutions; revives patients who fainted
Simultaneously with Lewis, a Soviet chemist Mikhail Usanovich from Tashkent, developed a general theory that does not restrict acidity to hydrogen-containing compounds, but his approach, published in 1938, was even more general than Lewis theory. Usanovich's theory can be summarized as defining an acid as anything that accepts negative species or donates positive ones, and a base as the reverse. This pushed the concept of acid-base reactions to its logical limits, and even redefined the concept of redox (oxidation-reduction) as a special case of acid-base reactions, and so did not become wide spread, despite being easier to understand than Lewis theory, which required detailed familiarity with atomic structure. Some examples of Usanovich acid-base reactions include:
This acid-base theory was a revival of oxygen theory of acids and bases, proposed by German chemist Hermann Lux in 1939, further improved by Håkon Flood, circa 1947 and is still used in modern geochemistry and electrochemistry of molten salts. This definition describes an acid as an oxide ion (O2−) acceptor and a base as an oxide ion donor. For example,
In 1963, Ralph Pearson proposed an advanced qualitative concept known as Hard Soft Acid Base principle, later made quantitative with help of Robert Parr in 1984. 'Hard' applies to species that are small, have high charge states, and are weakly polarizable. 'Soft' applies to species that are large, have low charge states and are strongly polarizable. Acids and bases interact, and the most stable interactions are hard-hard and soft-soft. This theory has found use in organic and inorganic chemistry.
Augustus Svante Arrhenius proposed in his study of electrolytic dissociation that in an aqueous solution, a strong electrolyte only exists in the form of ions, whereas a weak electrolyte exists partly as ions and partly as molecules. Years later, specifically three, Arrhenius extended this theory by suggesting that acids are neutral compounds that ionize when dissolved in water to give hydrogen ions and a corresponding negative ion. Bases are neutral compounds that either dissociate or ionize in water to give hydroxide ions and a positive ion.
An Arrhenius Acid, therefore, is any substance that ionizes when it dissolves in water to give hydrogen ion(s). For example,
HCl yields to H+ (aq) + Cl-(aq)
An Arrhenius Base is any substance that gives hydroxide ions when it dissolves in water. For example,
NaOH yields to Na+(aq) + OH-(aq)
The Arrhenius Theory has several disadvantages. It is only applicable for aqueous solutions. The theory also does not explain why not all compounds containing a proton give an acidic solution and why substances not containing hydroxides are able to produce basic solutions.
Bronsted-Lowry Theory
One change suggested by Johannes Bronsted and Thomas Lowry was the existence of a hydronium ion. They proposed that hydrogen ions cannot exist in water but rather covalently bond with water. They based their definitions of acids and bases on the behavior of hydrogen ions. Acids are substances from which a proton can be removed. bases are substances that binds protons. It can be simplified as Bronsted-Lowry acids being proton donors and Bronsted-Lowry bases as proton acceptors. Take note that protons are not actually being given but rather the bond between the proton and another element is being broken to form a new bond to the base. This new definition broadens the range of acids and bases but still requires a protic acid. The concept of conjugates is an additional concept introduced by this theory. A conjugate base is the species remaining when a proton is removed from an acid. A conjugate acid is the species formed when a proton is transferred to a base. For example,
Lewis Theory
The theory proposed by Lewis further widens the classification of acids and bases to non-aqueous solutions. He proposed that the electrons and not the protons are the ones transferring. The acids and bases can be charged or unionized. A Lewis acid is a species that can accept a pair of electrons while a Lewis base is a species that can donate a pair of electrons. Note that these species can be ions or molecules. For example,
Solvent System Theory
One of the limitations of Arrhenius definition was its reliance on water solutions. E. C. Franklin studied the acid-base reactions in liquid ammonia in 1905 and pointed out the similarities to water-based Arrhenius theory, and Albert F. O. Germann, working with liquid COCl2, generalized Arrhenius definition to cover aprotic solvents and formulated the solvent system theory in 1925.
Germann pointed out that in many solvents there is a certain concentration of a positive species, solvoniumsolvate (earlier lyate) anions, in equilibrium with the neutral solvent molecules. For example, water and ammonia undergo such dissociation into hydronium and hydroxide, and ammonium and amide, respectively: (earlier lyonium) cations and negative species,
2 H2OH3O+ + OH−
2 NH3NH+ 4 + NH− 2
Some aprotic systems also undergo such dissociation, such as dinitrogen tetroxide into nitrosonium and nitrate, antimony trichloride into dichloroantimonium and tetrachloroantimonate, and phosgene into chlorocarboxonium and chloride.
N2O4NO+ + NO− 3
2 SbCl3SbCl+ 2 + SbCl− 4
COCl2COCl+ + Cl−
A solute causing an increase in the concentration of the solvonium ions and a decrease in the solvate ions is defined as an acid and one causing the reverse is defined as a base. Thus, in liquid ammonia, KNH2NH−
2) is a strong base, and NH4NO3 (supplying NH+
4) is a strong acid. In liquid sulfur dioxide (SO2), thionyl compounds (supplying SO2+) behave as acids, and sulfites (supplying SO2−
3) behave as bases. (supplying (
The non-aqueous acid-base reactions in liquid ammonia are similar to the reactions in water:
Since solvent-system definition depends on the solvent as well as on the compound itself, the same compound can change its role depending on the choice of the solvent. Thus, HClO4 is a strong acid in water, a weak acid in acetic acid, and a weak base in fluorosulfonic acid. This was seen as both a strength and a weakness, since some substances, such as SO3 and NH3 were felt to be acidic or basic on their own right. On the other hand, solvent system theory was criticized as too general to be useful; it was felt that there was something intrinsically acidic about hydrogen compounds, not shared by non-hydrogenic solvonium salts.
The first scientific concept of acids and bases was provided by Antoine Lavoisier, circa 1776. Since Lavoisier's knowledge of strong acids was mainly restricted to oxoacids, such as HNO3 (nitric acid) and H2SO4 (sulfuric acid), which tend to contain central atoms in high oxidation states surrounded by oxygen, and since he was not aware of the true composition of the hydrohalic acids (HF, HCl, HBr, and HI), he defined acids in terms of their containing oxygen, which in fact he named from Greek words meaning "acid-former" (from the Greek oxys meaning "acid" or "sharp" and geinomai meaning "engender"). The Lavoisier definition was held as absolute truth for over 30 years, until the 1810 article and subsequent lectures by Sir Humphry Davy in which he proved the lack of oxygen in H2S, H2Te, and the hydrohalic acids. However, Davy failed to develop a new theory, concluding that "acidity does not depend upon any particular elementary substance, but upon peculiar arrangement of various substances". One notable modification of oxygen theory was provided by Berzelius, who stated that acids are oxides of nonmetals while bases are oxides of metals.
Hydrogen Theory of Acids
This definition was proposed by Justus von Liebig, circa 1838, based on his extensive works on the chemical composition of organic acids. This finished the doctrinal shift from oxygen-based acids to hydrogen-based acids, started by Davy. According to Liebig, an acid is a hydrogen-containing substance in which the hydrogen could be replaced by a metal. Liebig's definition, while completely empirical, remained in use for almost 50 years until the adoption of the Arrhenius definition.
NaOH H2O Na+(aq) + OH-(aq)
Ca(OH)2(s) H2O Ca2+(aq) + 2 OH-(aq)
H3O+ + OH−